This guide covers the building blocks of all substances—atoms, elements, and the Periodic Table. Mastering this is key to understanding all of chemistry.

1. Short Notes: Core Concepts

A. The Atom: The Basic Unit

• The atom consists of a central nucleus containing protons (+ charge) and neutrons (no charge).

• Electrons (- charge) orbit the nucleus in specific energy levels called shells (labeled K, L, M, N... or 1, 2, 3, 4...).

• Atomic Number (Z): This is the number of protons. It is the unique identity of an element. In a neutral atom, Protons = Electrons.

• Mass Number (A): The total number of Protons + Neutrons.

B. Electronic Configuration

• This is the arrangement of electrons in the shells. You must know this for the first 20 elements (up to Calcium).

• Shell Capacity: 1st shell (K) holds 2, 2nd (L) holds 8, 3rd (M) holds 8 (for O/L purposes up to Ca).

• Example: Magnesium (Mg)

  • Atomic Number = 12. So, it has 12 protons and 12 electrons.

  • Electronic Configuration: 2, 8, 2 (2 in the first shell, 8 in the second, 2 in the third).

C. Isotopes

• Atoms of the same element (same number of protons) but with a different number of neutrons.

• This means they have the same Atomic Number (Z) but a different Mass Number (A).

• Example: Carbon-12 (¹²C) has 6 protons, 6 neutrons. Carbon-14 (¹⁴C) has 6 protons, 8 neutrons.

D. The Modern Periodic Table

• A chart that organizes elements by their increasing atomic number.

• Periods (Horizontal Rows): The period number tells you the number of electron shells an atom has.

  • Example: Sodium (Na: 2, 8, 1) has 3 shells, so it's in Period 3.

• Groups (Vertical Columns): The group number tells you the number of valence electrons (electrons in the outermost shell). Elements in the same group have similar chemical properties.

  • Example: Chlorine (Cl: 2, 8, 7) has 7 valence electrons, so it's in Group VII.

• Noble Gases (Group VIII or 0): They have a full outer shell, making them very stable and unreactive.

E. Periodic Trends (How Properties Change)

• First Ionization Energy: Energy needed to remove the outermost electron.

  • Increases as you go across a period (left to right).

  • Decreases as you go down a group.

• Electronegativity: The ability of an atom to attract electrons in a chemical bond.

  • Increases as you go across a period.

  • Decreases as you go down a group.

F. Valency and Chemical Formulae

• Valency: The combining power of an element, determined by the number of electrons it needs to lose, gain, or share to become stable (like a noble gas).

  • Group I metals (e.g., Na) lose 1 electron -> Valency 1.

  • Group II metals (e.g., Mg) lose 2 electrons -> Valency 2.

  • Group VII non-metals (e.g., Cl) gain 1 electron -> Valency 1.

• Writing Formulas: Use the "criss-cross" method. The valency of one element becomes the subscript of the other.

  • Example: Calcium (valency 2) and Chlorine (valency 1) -> Ca₁Cl₂ -> CaCl₂

2. Tips & Tricks for the Exam

• Atomic Math:

  • Z = p (Atomic number = number of protons)

  • A = p + n (Mass number = protons + neutrons)

  • n = A - Z (Neutrons = Mass number - Atomic number)

• Periodic Table is your Cheat Sheet: If you know an element's position, you know its electron shells and valence electrons, which tells you its properties!

• "Down & Out": When going down a group, the outermost electron is further out. This makes it easier to remove (lower ionization energy) and harder for the nucleus to attract other electrons (lower electronegativity).

• Practice Formulas: The only way to get good at writing chemical formulas is to practice. Write down common ions (e.g., Sulphate SO₄²⁻, Nitrate NO₃⁻, Ammonium NH₄⁺) and practice combining them.

3. Important Points & Common Exam Questions

• Find the Element: A very common question gives you the atomic number and asks you to determine everything about it.

  • Example Question: An element has Z=19.

    1. Write its electronic configuration. (Answer: 2, 8, 8, 1)

    2. Identify its Group and Period. (Answer: Group I, Period 4)

    3. Is it a metal or non-metal? (Answer: Metal, as it's in Group I)

    4. What is the formula of its oxide? (Answer: It forms K⁺ ion, Oxygen forms O²⁻ ion -> K₂O)

• Explain the Trends: Be ready to explain why trends occur.

  • Why does electronegativity increase across Period 3?

  • Answer: Across the period, the number of protons in the nucleus increases, so the nuclear charge gets stronger. This pulls the electron shells closer and attracts bonding electrons more powerfully.

• Isotopes: You'll be asked to calculate the number of subatomic particles in a specific isotope, like ³⁵Cl₁₇.

  • Answer: Protons = 17, Electrons = 17, Neutrons = 35 - 17 = 18.

• Properties of Metals vs. Non-metals: Know the key differences (lustre, conductivity, type of oxide formed - basic for metals, acidic for non-metals).